Acids and Bases Some Definitions

Chapter 3 important inorganic compounds Acids and Bases Some Definitions • Arrhenius acids and bases Acid: Substance that, when dissolved in water, increases the concentration of hydrogen (hydronium ions) (protons, H+ or H3O+). Base: Substance that, when dissolved in water, increases the concentration of hydroxide ions. Cont…… • Brønsted–Lowry: must have both 1. an Acid: Proton donor and 2. a Base: Proton acceptor Brønsted-Lowry acids and bases are always paired. The Brønsted-Lowry acid donates a proton, while the Brønsted-Lowry base accepts it. Which is the acid and which is the base in each of these rxns? What Happens When an Acid Dissolves in Water? • Water acts as a Brønsted–Lowry base and abstracts a proton (H+) from the acid. • As a result, a hydronium ion (H3O+) is formed. Ionization of HCl and formation of hydronium ion, H3O+ H2O + Proton acceptor HCl Proton donor  H3O+ + Cl- A Brønsted–Lowry acid… …must have a removable (acidic) proton. HCl, H2O, H2SO4 A Brønsted–Lowry base… …must have a pair of nonbonding electrons. NH3, H2O Lewis defn of acid and base Lewis acid; a substance that accepts pair of electrons and capable of dative bond formation. Lewis bases; a substance donates pair of electrons and forms a dative bond. Example ; Properties of Acids  Acids taste sour  Acids affect indicators  Blue litmus turns red  Methyl orange turns red  Acids have a pH lower than 7  Acids are proton (hydrogen ion, H+) donors  Acids react with active metals, produce H2  Acids are electrolytes Acids reacts with carbonates and sulphites to produce salt oxide and water  Acids neutralize bases to produce a salt and water Acids change Indicators Blue litmus paper turns red in contact with an acid. Acids React with Active Metals Acids react with active metals to form salts and hydrogen gas. Mg + 2HCl  MgCl2 + H2(g) Zn + 2HCl  ZnCl2 + 4H2(g)  reaction of copper with hot and conc. H2SO4 and HNO3  Cu + 4HNO3  Cu + 2H2SO4 Cu(NO3)2 + 2NO2 + 2H2O CuSO4 + SO2 + 2H2O Reaction of acids with carbonates and sulphites • 2HCl + Na2CO3 Hydrochloric Sodium acid carbonate • HNO3 + Na2SO3 Nitric acid sodium sulphite 2NaCl + CO2 + H20 Sodium carbon chloride dioxide water NaNO3 + SO2 + H2O Sodium Sulfur nitrate dioxide Water Acids Neutralize Bases HCl + NaOH  NaCl + H2O Neutralization reactions ALWAYS produce a salt and water. Classification of acids Based on the relative amount of acid in a given solution  Concentrated- large acid small water- 70% HNO3, 99% CH3COOH  Diluted- small acid large water content – 2% HNO3 10% CH3COOH Based on degree of dissociation to ions  Strong – completely dissociate to ions- H2SO4, HCl, HNO3  Weak – partially or slightly dissociate to ions – H2CO3, H2SO3, H2C2O4 Concept Test • Which of the following "molecular" pictures best represents a concentrated solution of the weak acid HA? A B Strong Acids vs. Weak Acids Strong acids are assumed to be 100% ionized in solution (good proton donors). HCl H2SO4 HNO3 Weak acids are usually less than 5% ionized in solution (poor proton donors). H3PO4 HC2H3O2 Organic acids Strong Acid Dissociation Weak Acid Dissociation Based on number of ionizable hydrogen per a molecule Monoprotic acids Diprotic acids HCl H2SO4 HC2H3O2 H2CO3 Triprotic acids H3PO4 HNO3  based on composition  binary acids- contain two elements only HCl, HF, H2S Oxy acids or ternary acids – contain three elements HNO3, H2SO4,H3PO4 Preparation of acids • Reaction of acidic oxide with water  N2O5 + H2O 2HNO3 (aq)  SO2 + H2O H2SO3 (aq) • Direct combination of non-metals  H2 (g) + Cl2 (g) 2HCl (g) When HCl gas dissolved in water gives HCl (aq) acid • From reaction of non volatile acids with salts of volatile acids  NaNO3(s) + H2SO4(l) NaHSO4(s) + HNO3(l) Uses of acids • H3PO4 - soft drinks, fertilizer, detergents • H2SO4 – leading industrial chemical, fertilizer, car batteries • HNO3 – fertilizers, explosives (TNT and TNG), rubber, plastic, drugs • HCl - gastric juice ,pickling of iron, photographic films, drugs, plastics, PVC • HC2H3O2 - vinegar Acids you must know: Strong Acids Weak Acids Sulfuric acid, H2SO4 Phosphoric acid, H3PO4 Hydrochloric acid, HCl Acetic acid, HC2H3O2 Nitric acid, HNO3 Hydrobromic acid, HBr Hydroiodic acid, HI Perchloric acid, HClO4 Oxalic acid, H2C2O4 Hydrofluoric acid, HF Properties of Bases  Bases taste bitter  Bases affect indicators  Red litmus turns blue  Phenolphthalein turns purple  Bases have a pH greater than 7  Bases are proton (hydrogen ion, H+) acceptors  Solutions of bases feel slippery  Bases neutralize acids to produce a salt and water Strong and weak bases • Strong bases • Weak bases  Almost completely ionized in Partially ionized in aqueous solution aqueous solution  Are weak electrolytes  Group IA and IIA lower member bases are strong  Mg(OH)2, NH4OH bases  Are strong electrolytes NH4OH(aq) NH4+(aq)+ OH-(aq) KOH(aq) K+(aq) + OH-(aq) Concentrated and diluted bases • Concentrated base • Diluted base  Contain relatively large amount of base in a given volume of solution  Expressed in terms of mole per litter or molarity  Eg; 10M of NaOH, 10M NH4OH  Contain relatively small amount of base in a given volume of solution  Eg; 0.05M NaOH, 0.5M NH4OH Examples of Bases  Sodium hydroxide (lye), NaOH  Potassium hydroxide, KOH  Magnesium hydroxide, Mg(OH)2  Calcium hydroxide (lime), Ca(OH)2 Preparation of bases • Reaction of highly reactive metals( IA &IIA) with water 2Li (s) + 2H2O(l) 2LiOH(l) + H2 (g) Ca (s) + 2H2O (l) Ca(OH)2(aq) +H2(g) • Reaction of group IA & IIA oxides with water Na2O (s) + H2O(l) 2NaOH(aq) 2BaO(s) + 2H2O(l) 2Ba(OH)2 • Double displacement reaction Soluble salt reacts with soluble base to produce another soluble base and insoluble salt. Ba(OH)2 (aq) + K2SO4 (aq) 2KOH(aq) + BaSO4(s) Ca(OH)2(aq) + Na2CO3(aq) 2NaOH(aq) + CaCO3(s) Uses of Bases • • • • NaOH - lye, drain and oven cleaner Mg(OH)2 - laxative, antacid NH4OH - cleaners, fertilizer Ca(OH)2- to remove soil acidity -test CO2 -Cement production Naming Acids, Bases, Salts • Bases & Salts- Ionic Name the same as ionic compounds • Acids Binary • 2 elements, first hydrogen • Hydro___ic acid Ternary • 3 elements, change ending • -ate  ic, -ite  ous pH Scale • pH a measure of the concentration of H3O+ ions in solution measured with a pH meter or an indicator with a wide color range 14 0 7 INCREASING ACIDITY NEUTRAL INCREASING BASICITY Acids Have a pH less than 7 Bases have a pH greater than 7 pH Scale pH of Common Substances Try to determine the name of the chemical substance indicated in the pH scale. pH scale is logarithmic pH and pOH • Mathematically pH of a solution is calculated as the negative logarithm of concentration of hydrogen ion. pH = -log[H+] [H+] = antilog(-pH) pOH of a solution also calculated as the negative logarithm of concentration of hydronium ion. pOH = -log[OH-] [OH-] = antilog(-pOH) • The sum of pH and pOH in a given solution is 14 Ionic product of water 2H2O H3O+ + OH- Kw is the ion-product constant. Kw is the product of the molar concentrations of H3O+ and OH- ions at a particular temperature + Kw = [H 3O ][OH ] ________ [H2O] is constant 2 [H2O] Kw = [H3O+][OH-] = 1.0 × 10-14 at 25°C pKw = -log Kw [H3O+] = [OH-] Solution is neutral [H3O+] > [OH-] Solution is acidic [H3O+] < [OH-] Solution is basic • Can incorporate pH and pOH Kw = [H3O+][OH-] -log Kw = -log [H3O+][OH-] pKw = pH + pOH = 14 (at 25°C) Cont…. • Calculate the pH of a solution with hydrogen ion concentration of 1×10-5 M? pH = -log[H+] = -log(1×10-5) = -(-5 log(10) = 5×1 = 5 • Calculate the hydrogen ion concentration of an aqueous solution when its pH is 9. • Solution: pH = – log [H+] –pH = log [H+] [H+] = antilog (–pH) = 10–9 = 1.0 x 10–9 mol L–1 Cont…. • The Relationship Between pH and Concentration • pH scale is logarithmic, a solution of pH 1 has 10 times higher concentration of [H+] than that of a solution of pH 2 • Calculate the pH value of 0.05 M H2SO4. Assume that H2SO4 is completely ionized. • Solution: H2SO4 is a strong acid and ionizes as • H2SO4 (aq) → 2H+(aq) + SO42–(aq) 0.05 2 × 0.05 • Molarity of H+ = [H+] = 0.1M pH =-log(H+)= -log(0.1) = 1 Cont… • Find out the pH and pOH of a 0.0001M solution of NaOH? • Solution: NaOH is a strong base. It ionizes completely NaOH(aq) → Na+(aq) + OH– 0.0001M 0.0001M 0.0001M • Therefore, 0.0001M solution of NaOH produces 0.0001M OH ions [OH–] = 0.0001M= 1 x 10–4 • pOH = -log[OH–] = - log1 x 10–4 = 4 • pH = 14 - 4= 10 Cont…. • A salt is an ionic compound consisting of a cation (positive ion) and an anion (negative ion) . • It is formed when the replaceable hydrogen ions of an acid are completely or partially replaced by metal ions or ammonium ions • Formation Reactions: 1.Neutralization HCl (aq) + NaOH(aq) → NaCl(aq) + H2O(g) 2. Metal-Acid Reaction: Zn (s) + H2SO4 (aq) → ZnSO4 (aq) + H2 (g) 3. Carbonate-Acid Reaction Na2CO3 (aq) + 2HNO3 (aq) → NaNO3(aq) + H2O(l) + CO2(g) Cont… 4. Reaction of metal oxide with acid CuO(s) + H2SO4 (aq) → CuSO4(aq) + H2O(aq) ZnO (s) + 2HNO3 (aq) → Zn(NO3)2 (aq) + H2O (l) 5. Reaction of acidic oxide with a base Ca(OH)2 (s) + CO2(g) → CaCO3(s) + H2O (l) 6. Reaction of acidic oxide with basic oxide Na2O (s) + SO3(g) → Na2SO4 (s) 7. Formation of insoluble salts by double displacement Reactions of two soluble salts gives one soluble and one insoluble salts Na2CO3(aq) + CuSO4(aq) → CuCO3 (s) + Na2SO4 (aq) Classification of Salts • Salts are classified based on the extent of hydrogen replacement and the nature of the parent acid/base • Normal Salts: Formed by the complete replacement of all replaceable hydrogen atoms in an acid Examples: NaCl, Na2SO4 • Acidic Salts: Formed by the partial replacement of replaceable hydrogen atoms in a poly-protic acid. They still contain replaceable H and usually form acidic solutions. Example, NaHCO3 , NaHSO4 • Basic Salts: Formed by the partial neutralization of a poly-hydroxy base. They contain hydroxide ions. Example: Al(OH)2Cl, Zn(OH)NO3 Some important salts and their uses • NaCl : Common table salt, used for food preservation, and as a raw material for making chemicals like washing soda and baking soda. • CaCO3: Found in chalk, marble, and limestone; used in construction materials. • NH4NO3: A key component in artificial fertilizers. • NaHCO3: Baking soda; used as an antacid and in baking • FeSO4 : Iron tablet • BaSO4 : Barium meal- for gastrointestinal x-ray • KNO3: Gun powder, explosive, fertilizer • FeCl3 : Waste water treatment, etching printed circuits • CaSO4 .2H2O : Plastering walls, supporting fractured bone • CuSO4 : Electroplating, fungicide Properties of salts  Salts can be classified depending on the anion (negative ion) they possess, because the anion is partly responsible for the solubility of the salt. 1. Solubility of salts • When salts are added to water, Some salts are soluble in water; some are insoluble while some salts are only slightly soluble. 2. Tendency to absorb water from the atmosphere or release water to the atmosphere. • Salts can be classified as hygroscopic, deliquescent and efflorescent Hygroscopic salts are those which absorb water from the atmosphere but remain solid. • Example: Anhydrous copper (II) sulphate, CuSO4 Deliquescent salts absorb water from the atmosphere to form a solution. • The process of absorbing water from the atmosphere by a solid to form a solution is called deliquescence. Example: Calcium chloride(CaCl2), Sodium nitrate (NaNO3), Iron(III) chloride (FeCl3) Cont… • Efflorescent salts lose their water of crystallization to the atmosphere. • The loss of water of crystallization by solid crystals to the atmosphere is known as efflorescence. • It is very important to note that all deliquescent substances are hygroscopic, but all hygroscopic substances are not necessarily deliquescent. • Example Hydrated sodium carbonate (Na2CO3•10H2O), Hydrated sodium sulphate (Na2SO4•10H2O) Cont…. 3. Aqueous solutions of soluble salts are electrolytes, because they release mobile positive and negative ions in solution. • NaNO3 (aq) → Na+ (aq) + NO3– (aq) • CaCl2 (aq) → Ca2+(aq) + 2Cl– (aq) 4. Thermal stability of salts: • Thermal stability of a salt is the property of a salt to resist irreversible change in its chemical or physical structure, often by resisting decomposition at a high relative temperature. • Different salts containing the same anion are heated, they may not show similar behavior. Some salts are thermally stable while others undergo decomposition. Cont…. A. Thermal decomposition of carbonates • All the carbonates of Group IIA undergo thermal decomposition to the metal oxide and carbon dioxide gas. • Carbonates are salts containing a carbonate, CO32-, as an anion. • If “M” represents any group IIA metal, the following describes the decomposition: MCO3(s) → MO(s) + CO2 • The carbonates of Group IIA are: MgCO3, CaCO3, SrCO3, BaCO3 • The carbonates become more thermally stable down the group. In Group IA, lithium carbonate behaves in the same way, producing lithium oxide and carbon dioxide: Li2CO3(s) → Li2O(s)+CO2 • The rest of the Group IA carbonates do not decompose at laboratory temperatures Cont… • Most transition metals decompose on heating to give the metal oxides and carbon dioxide: • Example CuCO3 (s) → CuO (s) + CO2 (g) B. Thermal decomposition of nitrates • Nitrates are salts containing a nitrate, NO3–, as an anion. The Group IIA and most transition metal nitrates undergo thermal decomposition to the metal oxide, nitrogen dioxide and oxygen gas. Nitrate → Metal oxide + Nitrogen dioxide + Oxygen • Example 2Mg(NO3)2 (s) → 2MgO (s) + 4NO2 (g) + O2 (g) 2Pb(NO3)2 (s) → 2PbO (s) + 4NO2 (g) + O2 (g) Cont… • Nitrates of sodium and potassium decompose on heating to give nitrites (instead of the oxides) and oxygen gas; 2NaNO3 (s) → 2NaNO2 (s) + O2 (g) 2KNO3 (s) → 2KNO2 (s) + O2 (g) • The nitrate of lithium differ from those of sodium and potassium; they decompose on heating in the following manner: 4LiNO3 (s) → 2Li2O (s) + 4NO2 (s) + O2 (g) • Note that both the nitrates and carbonates of lithium are thermally unstable similar to Group IIA. Cont… • Flame tests are used to identify the presence of a relatively small number of metal ions in a compound. • Certain metals give a characteristic color to a Bunsen flame when their solid salts or moist salts are heated directly in the flame. • A flame test is commonly used to identify the presence of lithium, sodium, potassium, calcium, strontium and barium ions in salts. Task ; write the colors of some ions from a demonstration in lab.